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Ch.18 - Aqueous Ionic Equilibrium
Tro - Chemistry: A Molecular Approach 6th Edition
Tro6th EditionChemistry: A Molecular ApproachISBN: 9780137832217Not the one you use?Change textbook
All textbooksTro 6th EditionCh.18 - Aqueous Ionic EquilibriumProblem 54
Chapter 18, Problem 54

A 100.0-mL buffer solution is 0.100 M in NH3 and 0.125 M in NH4Br. What mass of HCl can this buffer neutralize before the pH falls below 9.00?

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1
First, we need to calculate the pOH from the given pH. We can use the formula pOH = 14 - pH.
Next, we need to calculate the concentration of OH- ions using the formula [OH-] = 10^-pOH.
Then, we need to calculate the concentration of NH3 in the buffer solution. We know that the concentration of NH3 is 0.100 M.
Now, we need to calculate the concentration of NH4+ ions in the buffer solution. We know that the concentration of NH4+ is 0.125 M.
Finally, we can calculate the mass of HCl that the buffer can neutralize. We know that the reaction between NH3 and HCl produces NH4+ and Cl-. Therefore, the amount of HCl that can be neutralized is equal to the amount of NH3 present in the buffer. We can use the formula mass = volume x concentration x molar mass to calculate the mass of HCl.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Buffer Solutions

Buffer solutions are mixtures that resist changes in pH when small amounts of acid or base are added. They typically consist of a weak acid and its conjugate base or a weak base and its conjugate acid. In this case, the NH3/NH4Br system acts as a buffer, maintaining the pH around 9.00 despite the addition of HCl.
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Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation relates the pH of a buffer solution to the concentration of its acidic and basic components. It is expressed as pH = pKa + log([A-]/[HA]). For the NH3/NH4+ buffer, knowing the pKa allows us to calculate how much HCl can be added before the pH drops below a specified value.
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Neutralization Reaction

A neutralization reaction occurs when an acid reacts with a base to form water and a salt. In this scenario, HCl (strong acid) will react with NH3 (weak base) to form NH4Cl and water. The extent of this reaction determines how much HCl can be added to the buffer before the pH changes significantly.
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Related Practice
Textbook Question

A 100.0-mL buffer solution is 0.175 M in HClO and 0.150 M in NaClO. c. What is the pH after addition of 85.0 mg of NaOH?

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. d. 105.0 mL of 0.12 M CH3NH2 ; 110.0 mL of 0.15 M CH3NH3Cl

Textbook Question

For each solution, calculate the initial and final pH after adding 0.010 mol of NaOH. a. 250.0 mL of pure water b. 250.0 mL of a buffer solution that is 0.195 M in HCHO2 and 0.275 M in KCHO2 c. 250.0 mL of a buffer solution that is 0.255 M in CH3CH2NH2 and 0.235 M in CH3CH2NH3Cl

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. a. 155.0 mL of 0.15 M NH3 ; 175.0 mL of 0.17 M HCl

Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. c. 55.0 mL of 0.15 M HF; 85.0 mL of 0.10 M NaF